1
Chapter Eight
Chemical Bonds
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Chapter Nine
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Chemical Bonds: A Preview
• Forces called chemical bonds hold atoms together
in molecules and keep ions in place in solid ionic
compounds.
• Chemical bonds are electrostatic forces; they
reflect a balance in the forces of attraction and
repulsion between electrically charged particles.
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Electrostatic Attractions and
Repulsions
Nuclei attract
electron(s)
Electrons
repel other
electrons
Nuclei repel
other nuclei
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Both attractions and
repulsions occur;
what’s the net
effect?? Answer: it
depends on the
distance between
nuclei …
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Chapter Nine
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Energy of Interaction
At 74 pm, attractive
forces are at a
maximum, energy is
at a minimum.
Closer together;
attraction increases.
Closer than 74
pm, repulsion
increases.
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Hydrogen nuclei far
apart; just a little
attraction.
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Chapter Nine
5
The Lewis Theory of
Chemical Bonding: An Overview
• Valence electrons play a fundamental role in chemical
bonding.
• When metals and nonmetals combine, valence electrons
usually are transferred from the metal to the nonmetal
atoms, giving rise to ionic bonds.
• In combinations involving only nonmetals, one or more
pairs of valence electrons are shared between the bonded
atoms, producing covalent bonds.
• In losing, gaining, or sharing electrons to form chemical
bonds, atoms tend to acquire the electron configurations of
noble gases.
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Chapter Nine
6
Lewis Symbols
• In a Lewis symbol, the chemical symbol for the element
represents the nucleus and core electrons of the atom.
• Dots around the symbol represent the valence electrons.
• In writing Lewis symbols, the first four dots are placed
singly on each of the four sides of the chemical symbol.
• Dots are paired as the next four are added.
• Lewis symbols are used primarily for those elements that
acquire noble-gas configurations when they form bonds.
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Chapter Nine
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Example 9.1
Give Lewis symbols for magnesium, silicon, and
phosphorus.
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Ionic Bonds and Ionic Crystals
• When atoms lose or gain electrons, they may
acquire a noble gas configuration, but do not
become noble gases.
• Because the two ions formed in a reaction between
a metal and a nonmetal have opposite charges,
they are strongly attracted to one another and form
an ion pair.
• The net attractive electrostatic forces that hold the
cations and anions together are ionic bonds.
• The highly ordered solid collection of ions is
called an ionic crystal.
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Chapter Nine
9
Formation of a Crystal of Sodium Chloride
Na donates an
electron to Cl …
… and
opposites
attract.
Sodium
reacts
violently in
chlorine gas.
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Chapter Nine
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Using Lewis Symbols
to Represent Ionic Bonding
• Lewis symbols can be used to represent ionic
bonding between nonmetals and: the s-block
metals, some p-block metals, and a few d-block
metals.
• Instead of using complete electron configurations
to represent the loss and gain of electrons, Lewis
symbols can be used.
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Chapter Nine
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Example 9.2
Use Lewis symbols to show the formation of ionic
bonds between magnesium and nitrogen. What are the
name and formula of the compound that results?
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Chapter Nine
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Energy Changes in Ionic
Compound Formation
Na(g)  Na+(g) + e–
I1 = +496 kJ/mol
Cl(g) + e–  Cl–(g)
EA1 = –349 kJ/mol
• From the data above, it doesn’t appear that the formation of
NaCl from its elements is energetically favored. However …
• … the enthalpy of formation of the ionic compound is more
important than either the first ionization energy or electron
affinity.
• The overall enthalpy change can be calculated using a stepwise procedure called the Born–Haber cycle.
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Chapter Nine
13
Energy Changes in Ionic Compound Formation (cont’d)
• The Born–Haber cycle
is a hypothetical
process, in which ΔHf
is represented by
several steps.
• What law can be used
to find an enthalpy
change that occurs in
steps??
∆ Hf° for NaCl is very
negative because …
… ∆ H5—the lattice
energy—is very
negative.
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Chapter Nine
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Example 9.3
Use the following data to determine the lattice energy of
MgF2(s): enthalpy of sublimation of magnesium, +146
kJ/mol; I1 for Mg, +738 kJ/mol; I2 for Mg, +1451 kJ/mol;
bond-dissociation energy of F2(g), +159 kJ/mol F2;
electron affinity of F, –328 kJ/mol F; enthalpy of
formation of MgF2(s), –1124 kJ/mol.
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Chapter Nine
15
Lewis Structures of
Simple Molecules
• A Lewis structure is a combination of Lewis
symbols that represents the formation of covalent
bonds between atoms.
• In most cases, a Lewis structure shows the bonded
atoms with the electron configuration of a noble
gas; that is, the atoms obey the octet rule. (H
obeys the duet rule.)
• The shared electrons can be counted for each atom
that shares them, so each atom may have a noble
gas configuration.
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Chapter Nine
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Lewis Structures (cont’d)
• The shared pairs of electrons in a molecule are called
bonding pairs.
• In common practice, the bonding pair is represented by a
dash (—).
• The other electron pairs, which are not shared, are called
nonbonding pairs, or lone pairs.
Each chlorine atom sees
an octet of electrons.
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Chapter Nine
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Some Illustrative Compounds
• Note that the two-dimensional Lewis structures do not necessarily show the correct
shapes of the three-dimensional molecules. Nor are they intended to do so.
• The Lewis structure for water may be drawn with all three atoms in a line: H–O–H.
• We will learn how to predict shapes of molecules in Chapter 10.
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Chapter Nine
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Multiple Covalent Bonds
• The covalent bond in which one pair of electrons is shared
is called a single bond.
In a double bond two pairs
• Multiple bonds can also form:
of electrons are shared.
In a triple bond three pairs
of electrons are shared.
Note that each atom obeys the octet
rule, even with multiple bonds.
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Chapter Nine
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The Importance of Experimental Evidence
• The Lewis structure
commonly drawn for oxygen
is
• But oxygen is paramagnetic,
and therefore must have
unpaired electrons.
• Lewis structures are a useful
tool, but they do not always
represent molecules correctly,
even when the Lewis
structure is plausible.
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Chapter Nine
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Polar Covalent Bonds and
Electronegativity
• Electronegativity (EN) is a measure of the ability of an
atom to attract its bonding electrons to itself.
• EN is related to ionization energy and electron affinity.
• The greater the EN of an atom in a molecule, the more
strongly the atom attracts the electrons in a covalent bond.
Electronegativity generally
increases from left to right
within a period, and it generally
increases from the bottom to the
top within a group.
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Chapter Nine
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Pauling’s Electronegativities
It would be a good idea to remember
the four elements of highest
electronegativity: N, O, F, Cl.
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Chapter Nine
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Example 9.4
Referring only to the periodic table inside the front
cover, arrange the following sets of atoms in the
expected order of increasing electronegativity.
(a) Cl, Mg, Si
(b) As, N, Sb
(c) As, Se, Sb
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Electronegativity Difference
and Bond Type
• Identical atoms have the same electronegativity and share a
bonding electron pair equally. The bond is a nonpolar
covalent bond.
• When electronegativities differ significantly, electron pairs
are shared unequally.
• The electrons are drawn closer to the atom of higher
electronegativity; the bond is a polar covalent bond.
• With still larger differences in electronegativity, electrons
may be completely transferred from metal to nonmetal
atoms to form ionic bonds.
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Chapter Nine
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Electronegativity Difference and
Bond Type
No sharp cutoff between
ionic and covalent bonds.
Cs—F bonds are “so polar”
that we call the bonds ____.
C—H bonds are
virtually nonpolar.
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Chapter Nine
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Depicting Polar Covalent Bonds
In nonpolar bonds,
electrons are shared
equally.
Polar bonds are also depicted
by partial positive and partial
negative symbols …
Unequal sharing
in polar covalent
bonds.
Polar bonds are often depicted
using colors to show
electrostatic potential (blue =
positive, red = negative).
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… or with a cross-based
arrow pointing to the more
electronegative element.
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Chapter Nine
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Example 9.5
Use electronegativity values to arrange the following
bonds in order of increasing polarity:
Br—Cl, Cl—Cl, Cl—F, H—Cl, I—Cl
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Chapter Nine
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Writing Lewis Structures:
Skeletal Structures
• The skeletal structure shows the arrangement of
atoms.
• Hydrogen atoms are terminal atoms (bonded to
only one other atom).
• The central atom of a structure usually has the
lowest electronegativity.
• In oxoacids (HClO4, HNO3, etc.) hydrogen atoms
are usually bonded to oxygen atoms.
• Molecules and polyatomic ions usually have
compact, symmetrical structures.
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Writing Lewis Structures: A Method
1.
2.
3.
4.
5.
Determine the total number of valence electrons.
Write a plausible skeletal structure and connect the atoms
by single dashes (covalent bonds).
Place pairs of electrons as lone pairs around the terminal
atoms to give each terminal atom (except H) an octet.
Assign any remaining electrons as lone pairs around the
central atom.
If necessary (if there are not enough electrons), move one
or more lone pairs of electrons from a terminal atom to
form a multiple bond to the central atom.
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Example 9.6
Write the Lewis structure of nitrogen trifluoride, NF3.
Example 9.7
Write a plausible Lewis structure for phosgene, COCl2.
Example 9.8
Write a plausible Lewis structure for the chlorate ion,
ClO3–.
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Chapter Nine
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Formal Charge
• Formal charge is the difference between the number of
valence electrons in a free (uncombined) atom and the
number of electrons assigned to that atom when bonded to
other atoms in a Lewis structure.
• Formal charge is a hypothetical quantity; a useful tool.
• Usually, the most plausible Lewis structure is one with no
formal charges.
• When formal charges are required, they should be as small
as possible.
• Negative formal charges should appear on the most
electronegative atoms.
• Adjacent atoms in a structure should not carry formal
charges of the same sign.
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Chapter Nine
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Formal Charge Illustrated
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Example 9.9
In Example 9.7, we wrote a Lewis structure for the
molecule COCl2, shown here as structure (a). Show
that structure (a) is more plausible than (b) or (c).
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Chapter Nine
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Resonance: Delocalized Bonding
• When a molecule or ion can be represented by two or more
plausible Lewis structures that differ only in the distribution of
electrons, the true structure is a composite, or hybrid, of them.
• The different plausible structures are called resonance
structures.
• The actual molecule or ion that is a hybrid of the resonance
structures is called a resonance hybrid.
• Electrons that are part of the resonance hybrid are spread out
over several atoms and are referred to as being delocalized.
Three pairs of
electrons are
distributed among
two bonds.
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Example 9.10
Write three equivalent Lewis structures for the SO3
molecule that conform to the octet rule, and describe
how the resonance hybrid is related to the three
structures.
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Chapter Nine
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Molecules that Don’t Follow
the Octet Rule
• Molecules with an odd number of valence electrons have
at least one of them unpaired and are called free radicals.
• Some molecules have incomplete octets. These are usually
compounds of Be, B, or Al; they generally have some
unusual bonding characteristics, and are often quite
reactive.
• Some compounds have expanded valence shells, which
means that the central atom has more than eight electrons
around it.
• A central atom can have expanded valence if it is in the
third period or lower (i.e., S, Cl, P).
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Example 9.11
Write the Lewis structure for bromine pentafluoride,
BrF5.
Example 9.12
A Conceptual Example
Indicate the error in each of the following Lewis
structures. Replace each by a more acceptable
structure(s).
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Chapter Nine
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Bond Order and Bond Length
• Bond order is the number of
shared electron pairs in a bond.
• A single bond has BO = 1, a
double bond has BO = 2, etc.
• Bond length is the distance
between the nuclei of two atoms
joined by a covalent bond.
• Bond length depends on the
particular atoms in the bond and
on the bond order.
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Example 9.13
Estimate the length of (a) the nitrogen-to-nitrogen bond
in N2H4 and (b) the bond in BrCl.
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Chapter Nine
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Bond Energy
• Bond-dissociation energy (D) is the energy required to
break one mole of a particular type of covalent bond in a
gas-phase compound.
• Energies of some bonds can differ from compound to
compound, so we use an average bond energy.
The H—H bond
energy is precisely
known …
… while the O—H bond
energies for the two bonds
in H2O are different.
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Trends in Bond Lengths and Energies
• The higher the order (for a particular type of
bond), the shorter and the stronger (higher
energy) the bond.
• A N=N double bond is shorter and stronger than a
N–N single bond.
• There are four electrons between the two positive
nuclei in N=N. This produces more electrostatic
attraction than the two electrons between the
nuclei in N–N.
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Calculations Involving Bond Energies
For the reaction N2(g) + 2 H2(g)  N2H4(g)
When the bonds of the
product form, 163 kJ plus
4(389 kJ) of energy is
liberated.
… we must supply
946 kJ …
… plus 2(436 kJ),
to break bonds of
the reactants.
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to occur …
∆ H = (+946 kJ) + 2(+436 kJ) + (–163 kJ) + 4(–389 kJ)
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Chapter Nine
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Example 9.14
Use bond energies from Table 9.1 to estimate the
enthalpy of formation of gaseous hydrazine. Compare
the result with the value of ∆ H°f [N2H4(g)] from
Appendix C.
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Chapter Nine
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Alkenes and Alkynes
• Hydrocarbons with double or triple bonds between carbon
atoms are called unsaturated hydrocarbons.
• Alkenes are hydrocarbons with one or more C=C double
bonds.
• Simple alkenes have just one double bond in their
molecules.
• The simplest alkene is C2H4, ethene (ethylene).
• Alkynes are hydrocarbons that have one or more carbon–
carbon triple bonds.
• The simplest alkyne is C2H2, ethyne (acetylene).
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Chapter Nine
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Molecular Models of Ethene and Ethyne
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Polymers
• Polymers are compounds in which many identical
molecules have been joined together.
• Monomers are the simple molecules which join together to
form polymers.
• Often, the monomers have double or triple bonds.
• The process of these molecules joining together is called
polymerization.
• Many everyday products and many biological compounds
are polymers.
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Chapter Nine
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Formation of Polyethylene
Another ethylene
molecule adds to a long
chain formed from
more ethylene
molecules.
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Chapter Nine
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Cumulative Example
Use data from this chapter and elsewhere in the text to
estimate a value for the standard enthalpy of formation
at 25 °C of HNCO(g). Assess the most likely source of
error in this estimation.
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Chapter Nine